electrochemistry

• Oxidation numbers, also known as oxidation state, is a number given to each atom or ion in a compound to keep track of how many electrons they have.
• In a single ion or molecular ion, the oxidation number tells us how many electrons have been lost or gained.
• Positive oxidation number indicates loss of electrons.
• Negative oxidation number indicates gain of electrons.
• Potassium sulfate is a salt of the common acid, so they are named without including the oxidation number of the non-metal.
• Magnesium nitrate is a salt of the common acid, so they are named without including the oxidation number of the non-metal.
• The oxidation number of N in NH4NO3 is +4.
• The oxidation number (ox.no.) refers to a single atom in a compound.
• Oxidation number rules dictate that the oxidation number (ox.no.) refers to a single atom in a compound.
• Oxidation numbers can be used to balance chemical equations.
• Roman numerals between brackets are used to show the oxidation number of an atom that can have multiple oxidation states, such as Fe(III) = iron with oxidation number +3.
• Writing overall redox reactions involves identifying the atoms which change in oxidation number, deducing the oxidation number changes, balancing the oxidation number changes, balancing the charges, and balancing the atoms.
• A reducing agent is a substance that reduces another atom or ion by causing it to gain electrons.
• An oxidising agent itself gets reduced, gains electrons, and therefore, the oxidation number of the oxidising agent decreases.
• Fe in FeO can be written as Iron(II) oxide.
• An oxidising agent is a substance that oxidises another atom or ion by causing it to lose electrons.
• In equation B, Fe oxidises Mg(0) to Mg (+2) and is itself reduced from Fe (+2) to Fe(0).
• Some substances can act both as oxidising and reducing agents, their nature is dependent upon what they are reacting with and the reaction conditions.
• Iron can be both +2 and +3 so Roman numerals are used to distinguish between them.
• A reducing agent itself gets oxidised, loses/donates electrons, and therefore, the oxidation number of the reducing agent increases.
• Fe in Fe O can be written as Iron(III) oxide.
• For a reaction to be recognised as a redox reaction, there must be both an oxidising and reducing agent.
• Roman numerals are used to show the oxidation states of transition metals which can have more than one oxidation number.
• Oxidising agents are substances that oxidise other species, gain electrons and are themselves reduced.
• Oxidation is the gain of oxygen, for example, Cu + H O → CuO + H.
• A disproportionation reaction is a reaction in which the same species is both oxidised and reduced, for example, Cl2 + 2AgNO3 → 2ClNO2 + AgCl.
• Redox reactions are reactions in which oxidation and reduction take place together, for example, Cu + Mg → Mg + Cu (Cu has been reduced and Mg has been oxidised).
• Oxidation can involve the loss of electrons, for example, Cu + Mg → Mg + Cu.
• Reduct ion causes a decrease in oxidation number, for example, 2 2 3 2 2 2+ 2+ 2+ 2 2 3 2 2 2+ 2+.
• Reduction can involve the gain of electrons, for example, Cu + Mg → Mg + Cu (Cu has been reduced and Mg has been oxidised).
• Oxidation also involves the loss of a hydrogen, for example, 2NH + 3Br → N + 6HBr.
• Reduction also involves the gain of a hydrogen, for example, 2NH + 3Br → N + 6HBr.
• Oxidation causes an increase in oxidation number, for example, Cu + Mg → Mg + Cu (change in oxidation number of Mg is +2 thus Mg is oxidised).
• Reduction is the loss of oxygen, for example, Cu+ H O → 2CuO + H.